13 review

Gibbs-Helmholtz Equation

• Gibbs free energy (G) and its relationship to enthalpy (H) and entropy (S).

  • ( G = H - T imes S )

  • Enthalpy (( \Delta H )) is negative for exothermic processes and positive for endothermic processes.

  • Entropy (( \Delta S )):

    • Increases during exothermic processes

    • Decreases during endothermic processes.

Solvation and Thermodynamics

• Solvent-solvent interactions are generally small.

• Heat of solution is determined by the balance of crystal lattice energy and solvation energy:

  • ( \Delta H_{solution} = \Delta H_{lattice} + \Delta H_{solvation} )

States of Matter & Stoichiometry

• During a process at constant volume, enthalpy equals heat observed.

• Entropy reflects energy distribution in a system:

  • In a solid state, molecules are more ordered than in a solution.

  • Dissolution processes increase disorder, thus increasing entropy.

Solubility Rules

• Always Soluble:

  • Nitrates (NO3-), Acetates (C2H3O2-), Group 1 ions (Li+, Na+, K+, etc.), Ammonium (NH4+), and Sulfates (SO4^2-).

  • Exceptions exist for some sulfates (e.g., Pb2+, Ag+).

Enthalpy and Entropy Relationships

• Reactions are favored based on the criteria of temperature and changes in enthalpy/entropy:

  • Negative enthalpy (( \Delta H < 0 )), positive entropy (( \Delta S > 0 )): Product favored at all temperatures.

  • Positive enthalpy (( \Delta H > 0 )), negative entropy (( \Delta S < 0 )): Reactant favored at all temperatures.

Types of Intermolecular Forces (IMFs)

• Ion-ion interactions: Attraction between ions of opposite charges.

• Dipole-dipole interactions: Between molecules with permanent dipoles.

• Hydrogen bonding: A type of dipole-dipole interaction involving H and highly electronegative atoms (O, N, F).

• London dispersion forces: Weakest forces present in all molecules, especially significant in nonpolar substances.

Crystal Lattice Energy

• Crystal lattice energy considered endothermic during dissolution.

• Smaller cations lead to larger lattice energies, as attractive forces increase with smaller ion size.

  • Example: LiCl has larger lattice energy compared to KCl due to the smaller size of Li+.

Factors Affecting Solubility

• The rate of dissolution is affected by the surface area of solids: larger surface area leads to faster dissolution.

• Colligative properties depend on the number of solute particles:

  • Vapor pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure.

Units of Concentration

• Percent by weight = (mass of solute / mass of solution) * 100

• Molarity (M) = moles of solute / liters of solution

• Molality (m) = moles of solute / kilograms of solvent

• Mole fraction (x) = moles of solute / total moles of solution.

Colligative Properties and Raoult's Law

• The vapor pressure of a solution is lower than that of a pure solvent due to non-volatile solutes:

  • ( P_{solvent} = X_{solvent} \times P_{solvent}^0 )

• Osmosis: Spontaneous movement of solvent across a semipermeable membrane from lower to higher solute concentration.

Boiling and Freezing Points

• Boiling Point Elevation: ( T_b = K_b imes m imes i )

• Freezing Point Depression: ( T_f = K_f imes m imes i )

• Both properties are influenced by the number of solute particles in solution.